This document provides an overview of biochemistry and its relevance to medicine. It discusses the following key points: - Biochemistry is the study of chemical processes in living organisms, from molecules to cells. It is essential for understanding normal body functions and disease states. - The human body is composed of biomolecules like proteins, lipids, carbohydrates, nucleic acids, and water, which are organized in a hierarchical structure from atoms to cells. - Intermolecular forces like covalent bonds, ionic bonds, hydrogen bonds, hydrophobic interactions, and van der Waals forces contribute to molecular stability. - Metabolic processes involve the breakdown and synthesis of biomolecules, deriving energy through oxidation

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CHAPTER
1Biochemical
Perspective to Medicine
Chapter at a Glance
The reader will be able to answer questions on the following topics:
¾¾History of biochemistry
¾¾Ionic bonds
¾¾Hydrogen bonding
¾¾Hydrophobic interactions
¾¾Principles of thermodynamics
¾¾Donnan membrane equilibrium
Biochemistry is the language of biology. The tools for
research in all the branches of medical science are mainly
biochemical in nature. The study of biochemistry is
essential to understand basic functions of the body. This
study will give information regarding the functioning of
cells at the molecular level. How the food that we eat is
digested, absorbed, and used to make ingredients of the
body? How does the body derive energy for the normal
day to day work? How are the various metabolic processes
interrelated? What is the function of genes? What is the
molecular basis for immunological resistance against
invading organisms? Answer for such basic questions can
only be derived by a systematic study of biochemistry.
Modern day medical practice is highly dependent on
the laboratory analysis of body fluids, especially the blood.
The disease manifestations are reflected in the composition
of blood and other tissues. Hence, the demarcation of
abnormal from normal constituents of the body is another
aim of the study of biochemistry.
The word chemistry is derived from the Greek word "chemi" (the
black land), the ancient name of Egypt. Indian medical science, even from
ancient times, had identified the metabolic and genetic basis of diseases.
Charaka, the great master of Indian Medicine, in his treatise (circa 400
BC) observed that madhumeha (diabetes mellitus) is produced by the
alterations in the metabolism of carbohydrates and fats; the statement
still holds good.
Biochemistry has developed as an offshoot of organic chemistry,
and this branch was often referred as "physiological chemistry". The
term "Biochemistry" was coined by Neuberg in 1903 from Greek
words, bios (= life) and chymos (= juice). One of the earliest treatises in
biochemistry was the "Book of Organic Chemistry and its Applications
to Physiology and Pathology", published in 1842 by Justus von Liebig
Hippocrates
460–377 BC
Charaka
400 BC
Sushrutha
500 BC

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4 Textbook of Biochemistry
The large amount of data, especially with regard to single nucleotide
polymorphisms (SNPs) that are available, could be harnessed by
"Bioinformatics". Computers are already helping in drug designing
process. Studies on oncogenes have identified molecular mechanisms of
control of normal and abnormal cells. Medical practice is now depending
more on the science of Medical Biochemistry. With the help of Human
genome project (HGP) the sequences of whole human genes are now
available; it has already made great impact on medicine and related
health sciences.
(1803–73), who introduced the concept of metabolism. The "Textbook
of Physiological Chemistry" was published in 1877 by Felix Hoppe-
Seyler (1825–95), who was Professor of Physiological chemistry at
Strausbourge University, France. Some of the milestones in the develop-
ment of the science of biochemistry are given in Table 1.1.
The practice of medicine is both an art and a science. The word
“doctor” is derived from the Latin root, "docere", which means “to
teach”. Knowledge devoid of ethical background may sometimes be
disastrous! Hippocrates (460 BC to 377 BC), the father of modern
medicine articulated "the Oath”. About one century earlier, Sushrutha
(?500 BC), the great Indian surgeon, enunciated a code of conduct for
the medical practitioners, which is still valid. He proclaims: “You must
speak only truth; care for the good of all living beings; devote yourself to
the healing of the sick even if your life be lost by your work; be simply
clothed and drink no intoxicant; always seek to grow in knowledge; in
face of God, you can take upon yourself these vows.”
Biochemistry is perhaps the most rapidly developing discipline
in medicine. No wonder, the major share of Nobel prizes in medicine
has gone to research workers engaged in biochemistry. Thanks to the
advent of DNA recombinant technology, genes can now be transferred
from one person to another, so that many of the genetically determined
diseases are now amenable to gene therapy. Many genes, (e.g. human
insulin gene) have already been transferred to microorganisms for large
scale production of human insulin. Advances in genomics like RNA
interference for silencing of genes and creation of transgenic animals
by gene targeting of embryonic stem cells are opening up new vistas
in therapy of diseases like cancer and AIDS. It is hoped that in future,
the physician will be able to treat the patient, understanding his genetic
basis, so that very efficient "designer medicine" could cure the diseases.
TABLE 1.1: Milestones in history of Biochemistry
Scientists
Rouelle
Lavoisier
Wohler
Berzelius
Louis Pasteur
Edward Buchner
Fiske and Subbarao
Lohmann
Hans Krebs
Avery and Macleod
Lehninger
Watson and Crick
Nirenberg
Holley
Khorana
Paul Berg
Kary Mullis
Year
1773
1785
1828
1835
1860
1897
1926
1932
1937
1944
1950
1953
1961
1963
1965
1972
1985
1990
2000
2003
2012
Landmark discoveries
Isolated urea from urine
Oxidation of food stuffs
Synthesis of urea
Enzyme catalysis theory
Fermentation process
Extracted enzymes
Isolated ATP from muscle
Creatine phosphate
Citric acid cycle
DNA is genetic material
TCA cycle in mitochondria
Structure of DNA
Genetic code in mRNA
Sequenced gene for tRNA
Synthesized the gene
Recombinant DNA technology
Polymerase chain reaction
Human genome project started
Draft human genome
Human genome project completed
ENCyclopedia Of DNA Elements
BIOMOLECULES
More than 99% of the human body is composed of 6
elements, i.e. oxygen, carbon, hydrogen, nitrogen, calcium
and phosphorus. Human body is composed of about 60%
water, 15% proteins, 15% lipids, 2% carbohydrates and
8% minerals. Molecular structures in organisms are built
from 30 small precursors, sometimes called the alphabets
of biochemistry. These are 20 amino acids, 2 purines,
3 pyrimidines, sugars (glucose and ribose), palmitate,
glycerol and choline.
In living organisms, biomolecules are ordered into
a hierarchy of increasing molecular complexity. These
biomolecules are covalently linked to each other to form
macromolecules of the cell, e.g. glucose to glycogen,
amino acids to proteins, etc. Major complex biomolecules
are proteins, polysaccharides, lipids and nucleic acids. The
macromolecules associate with each other by noncovalent
forces to form supramolecular systems, e.g. ribosomes,
lipoproteins.
Lavoisier
1743–1794
Berzelius
1779–1848
Friedrich
Wohler
1800–1882
Justus von Liebig
1803–1873
ENCODE
Frederick
Donnan
1870–1956
Louis Pasteur
1822–1895
Johannes van Albert Lehninger
der Waals1917–1986
NP 1910,
1837–1923

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Chapter 1: Biochemical Perspective to Medicine
Finally at the highest level of organization in the
hierarchy of cell structure, various supramolecular
complexes are further assembled into cell organelle. In
prokaryotes (e.g. bacteria; Greek word "pro" = before;
karyon = nucleus), these macromolecules are seen in a
homogeneous matrix; but in eukaryotic cells (e.g. higher
organisms; Greek word "eu" = true), the cytoplasm
contains various subcellular organelles. Comparison of
prokaryotes and eukaryotes are shown in Table 1.2.
5
STUDY OF METABOLIC PROCESSES
Our food contains carbohydrates, fats and proteins as
principal ingredients. These macromolecules are to be
first broken down to small units; carbohydrates to mono-
saccharides and proteins to amino acids. This process is
taking place in the gastrointestinal tract and is called
digestion or primary metabolism. After absorption, the
small molecules are further broken down and oxidized
to carbon dioxide. In this process, NADH or FADH2 are
generated. This is named as secondary or intermediary
metabolism. Finally, these reducing equivalents enter the
electron transport chain in the mitochondria, where they
are oxidized to water; in this process energy is trapped as
ATP. This is termed tertiary metabolism. Metabolism is
the sum of all chemical changes of a compound inside the
body, which includes synthesis (anabolism) and breakdown
(catabolism). (Greek word, kata = down; ballein = change).
electrons from the outer most orbit of an electropositive
atom to the outermost orbit of an electronegative atom. This
transfer results in the formation of a ‘cation’ and an ‘anion’,
which get consequently bound by an ionic bond. Common
examples of such compounds include NaCl, KBr and NaF.
With regard to protein chemistry, positive charges are
produced by epsilon amino group of lysine, guanidium
group of arginine and imidazolium group of histidine.
Negative charges are provided by beta and gamma carboxyl
groups of aspartic acid and glutamic acid (Fig.1.3).
Hydrogen Bonds
These are formed by sharing of a hydrogen between two
electron donors. Hydrogen bonds result from electrostatic
Fig. 1.1: Covalent bond
STABILIZING FORCES IN MOLECULES
Covalent Bonds
Molecules are formed by sharing of electrons between
atoms (Fig. 1.1).
Ionic Bonds or Electrostatic Bonds
Ionic bonds result from the electrostatic attraction
between two ionized groups of opposite charges
(Fig.1.2). They are formed by transfer of one or more
TABLE 1.2: Bacterial and mammalian cells
Prokaryotic cellEukaryotic cell
SizeSmallLarge; 1000 to 10,000 times
Cell wallRigidMembrane of lipid bilayer
NucleusNot definedWell defined
Organelles NilSeveral; including mitochondria
and lysosomes
Fig. 1.2: Ionic bond
Fig. 1.3: Ionic bonds used in protein interactions

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6 Textbook of Biochemistry
Johannes van der Waals (1837–1923). He was awarded
Nobel prize in 1910. These are short range attractive
forces between chemical groups in contact. Van der Waals
interactions occur in all types of molecules, both polar and
non-polar. The energy of the van der Waals interaction
is about 1 kcal/mol and are unaffected by changes in
pH. This force will drastically reduce, when the distance
between atoms is increased. Although very weak, van der
Waals forces collectively contribute maximum towards the
stability of protein structure, especially in preserving the
non-polar interior structure of proteins.
attraction between an electronegative atom and a hydrogen
atom that is bonded covalently to a second electronegative
atom. Normally, a hydrogen atom forms a covalent bond
with only one other atom. However, a hydrogen atom co-
valently bonded to a donor atom, may form an additional
weak association, the hydrogen bond with an acceptor atom.
In biological systems, both donors and acceptors are usually
nitrogen or oxygen atoms, especially those atoms in amino
(NH2) and hydroxyl (OH) groups.
With regard to protein chemistry, hydrogen releasing
groups are –NH (imidazole, in dole, peptide); –OH (serine,
threonine) and –NH2 (arginine, lysine). Hydrogen accepting
groups are COO— (aspartic, glutamic) C=O (peptide); and S–S
(disulphide). The DNA structure is maintained by hydrogen
bonding between the purine and pyrimidine residues.
WATER: THE UNIVERSAL SOLVENT
Water constitutes about 70 to 80 percent of the weight of
most cells. The hydrogen atom in one water molecule is
attracted to a pair of electrons in the outer shell of an oxygen
atom in an adjacent molecule. The structure of liquid water
contains hydrogen-bonded networks (Fig. 1.5).
The crystal structure of ice depicts a tetrahedral
arrangement of water molecules. On melting, the molecules
get much closer and this results in the increase in density
of water. Hence, liquid water is denser than solid ice. This
also explains why ice floats on water.
Water molecules are in rapid motion, constantly making
and breaking hydrogen bonds with adjacent molecules.
As the temperature of water increases toward 100°C, the
kinetic energy of its molecules becomes greater than the
energy of the hydrogen bonds connecting them, and the
gaseous form of water appears. The unique properties of
water make it the most preferred medium for all cellular
reactions and interactions.
Hydrophobic Interactions
Non-polar groups have a tendency to associate with each other
in an aqueous environment; this is referred to as hydrophobic
interaction. These are formed by interactions between
nonpolar hydrophobic side chains by eliminating water
molecules. The force that causes hydrophobic molecules
or nonpolar portions of molecules to aggregate together
rather than to dissolve in water is called the ‘hydrophobic
bond’ (Fig.1.4). This serves to hold lipophilic side chains of
amino acids together. Thus non-polar molecules will have
minimum exposure to water molecules.
Van Der Waals Forces
These are very weak forces of attraction between all atoms,
due to oscillating dipoles, described by the Dutch physicist
Fig. 1.4: Hydrophobic interaction Fig. 1.5: Water molecules hydrogen bonded

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Chapter 1: Biochemical Perspective to Medicine
a. Water is a polar molecule. Molecules with polar bonds
that can easily form hydrogen bonds with water can
dissolve in water and are termed “hydrophilic”.
b. It has immense hydrogen bonding capacity both with
other molecules and also the adjacent water molecules.
This contributes to cohesiveness of water.
c. Water favors hydrophobic interactions and provides a
basis for metabolism of insoluble substances.
Water expands when it is cooled from 4° C to 0° C,
while normally liquids are expected to contract due to
cooling. As water is heated from 0° C to 4° C, the hydrogen
bonds begin to break. This results in a decrease in volume
or in other words, an increase in density. Hence, water
attains high density at 4° C. However, above 4° C the effect
of temperature predominates.
7
A closed system approaches a state of equilibrium.
Any system can spontaneously proceed from a state of low
probability (ordered state) to a state of high probability
(disordered state). The entropy of a system may decrease
with an increase in that of the surroundings. The second
law may be expressed in simple terms as Q = T × ∆S,
where Q is the heat absorbed, T is the absolute temperature
and ∆S is the change in entropy.
Gibb's Free Energy Concept
The term free energy is used to get an equation combining
the first and second laws of thermodynamics. Thus, ∆G =
∆H – T∆S, where ∆G is the change in free energy, ∆H is
the change in enthalpy or heat content of the system and ∆S
is the change in entropy. The term free energy denotes a
portion of the total energy change in a system that is
available for doing work.
For most biochemical reactions, it is seen that ∆H is
nearly equal to ∆E. So, ∆G = ∆E – T∆S. Hence, ∆G or
free energy of a system depends on the change in internal
energy and change in entropy of a system.
PRINCIPLES OF THERMODYNAMICS
Thermodynamics is concerned with the flow of heat and
it deals with the relationship between heat and work.
Bioenergetics, or biochemical thermodynamics, is the
study of the energy changes accompanying biochemical
reactions. Biological systems use chemical energy to
power living processes. Standard Free Energy Change
It is the free energy change under standard conditions. It is
designated as ∆G0. The standard conditions are defined for
biochemical reactions at a pH of 7 and 1 M concentration,
and differentiated by a priming sign ∆G0´. It is directly
related to the equilibrium constant. Actual free energy
changes depend on reactant and product.
Most of the reversible metabolic reactions are near
equilibrium reactions and therefore their ∆G is nearly zero.
The net rate of near equilibrium reactions are effectively
regulated by the relative concentration of substrates
and products. The metabolic reactions that function far
from equilibrium are irreversible. The velocities of these
reactions are altered by changes in enzyme activity. A
highly exergonic reaction is irreversible and goes to
completion. Such a reaction that is part of a metabolic
pathway, confers direction to the pathway and makes the
entire pathway irreversible.
Laws of thermodynamics have many applications in
biology and biochemistry, such as study of ATP hydrolysis,
membrane diffusion, enzyme catalysis as well as DNA
binding and protein stability. These laws have been used to
explain hypothesis of origin of life.
First Law of Thermodynamics
The total energy of a system, including its surroundings,
remains constant. Or, ∆E = Q – W, where Q is the heat
absorbed by the system and W is the work done. This is
also called the law of conservation of energy. If heat is
transformed into work, there is proportionality between
the work obtained and the heat dissipated. A system is an
object or a quantity of matter, chosen for observation. All
other parts of the universe, outside the boundary of the
system, are called the surrounding.
Second Law of Thermodynamics
The total entropy of a system must increase if a
process is to occur spontaneously. A reaction occurs
spontaneously if ∆E is negative, or if the entropy of the
system increases. Entropy (S) is a measure of the degree
of randomness or disorder of a system. Entropy becomes
maximum in a system as it approaches true equilibrium.
Enthalpy is the heat content of a system and entropy
is that fraction of enthalpy which is not available to do
useful work.

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8 Textbook of Biochemistry
In Figure 1.6, the left compartment contains NaR,
which will dissociate into Na+ and R¯. Then Na+ can diffuse
freely, but R¯ having high molecular weight cannot diffuse.
The right compartment contains NaCl, which dissociates
into Na+ and Cl¯, in which case, both ions can diffuse freely.
Thus, if a salt of NaR is placed in one side of a
membrane, at equilibrium
Na+ × R¯ × H+ × OH¯ = Na+ × OH¯ × H+
To convey the meaning of the mathematical values, a
hypothetical quantity of each of the ion is also incorporated
in brackets. Initially 5 molecules of NaR are added to the
left compartment and 10 molecules of NaCl in the right
compartment and both of them are ionized (Fig.1.6A).
When equilibrium is reached, the distributions of ions are
shown in Figure 1.6B. According to Donnan's equilibrium,
the products of diffusible electrolytes in both the
compartments will be equal, so that
[Na+] L × [Cl¯ ] L = [Na+] R × [Cl¯ ] R
If we substitute the actual numbers of ions, the formula
becomes
9 × 4 in left = 6 × 6 in right
Donnan's equation also states that the electrical
neutrality in each compartment should be maintained. In
other words the number of cations should be equal to the
number of anions, such that
In left: Na+= R¯+ Cl¯; substituting: 9 = 5 + 4
In right: Na+ = Cl¯; substituting: 6 = 6
The equation should also satisfy that the number
of sodium ions before and after the equilibrium are the
same; in our example, initial Na+ in the two compartments
together is 5 + 10 = 15; after equilibrium also, the value is
9 + 6 = 15. In the case of chloride ions, initial value is 10
and final value is also 4 + 6 = 10.
In summary, Donnan's equations satisfy the following
results:
1. The products of diffusible electrolytes in both
compartments are equal.
2. The electrical neutrality of each compartment is
maintained.
3. The total number of a particular type of ions before
and after the equilibrium is the same.
4. As a result, when there is non-diffusible anion on
one side of a membrane, the diffusible cations are
more, and diffusible anions are less, on that side.
Three Types of Reactions
A. A reaction can occur spontaneously when ∆G is
negative. Then the reaction is exergonic. If ∆G is of
great magnitude, the reaction goes to completion and
is essentially irreversible.
B. When ∆G is zero, the system is at equilibrium.
C. For reactions where ∆G is positive, an input of energy
is required to drive the reaction. The reaction is termed
as endergonic. (Examples are given in Chapter 5).
Similarly a reaction may be exothermic (∆H is negative),
isothermic (∆H is zero) or endothermic (∆H is positive).
Energetically unfavourable reaction may be driven
forward by coupling it with a favourable reaction.
Glucose + Pi → Glucose-6-phosphate(reaction1)
ATP + H2O → ADP + Pi(reaction 2)
Glucose + ATP→ Glucose-6-phosphate+ADP (3)
Reaction 1 cannot proceed spontaneously. But the
2nd reaction is coupled in the body, so that the reaction
becomes possible. For the first reaction, ∆G0 is +13.8 kJ/
mole; for the second reaction, ∆G0 is –30.5 kJ/mole. When
the two reactions are coupled in the reaction 3, the ∆G0
becomes –16.7 kJ/mole, and hence the reaction becomes
possible. Details on ATP and other high-energy phosphate
bonds are described in Chapter 20.
Reactions of catabolic pathways (degradation or
oxidation of fuel molecules) are usually exergonic. On the
other hand, anabolic pathways (synthetic reactions or building
up of compounds) are endergonic. Metabolism constitutes
anabolic and catabolic processes that are well co-ordinated.
DONNAN MEMBRANE EQUILIBRIUM
When two solutions are separated by a membrane
permeable to both water and small ions, but when one of
the compartments contains impermeable ions like proteins,
distribution of permeable ions occurs according to the
calculations of Donnan.
A B
Fig. 1.6: Donnan membrane equilibrium

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Chapter 1: Biochemical Perspective to Medicine9
Clinical Applications of the Equation
1. The total concentration of solutes in plasma will be
more than that of a solution of same ionic strength
containing only diffusible ions; this provides the net
osmotic gradient (see under Albumin, in Chapter 28).
2. The lower pH values within tissue cells than in the
surrounding fluids are partly due to the concentrations
of negative protein ions within the cells being higher
than in surrounding fluids.
3. The pH within red cells is lower than that of the
surrounding plasma is due, in part, to the very high
concentration of negative non-diffusible hemoglobin
ions. This will cause unequal distribution of H+ ions
with a higher concentration within the cell.
4. The chloride shift in erythrocytes as well as higher
concentration of chloride in CSF are also due to
Donnan's effect.
5. Osmolarity of body fluid compartments and
sodium
concentration will follow Donnan equation.
6. Different steps of water purification
employ the
same principle and may be cited as an example of
industrial application of the equation.